Understanding Lewis

Structures
Dr. Mehr Nigar
 The Bond Length

ΔEd

Energy of interaction = nuclear-nuclear + electron-nuclear +

electron-electron attraction
repulsion repulsion

Bond Dissociation Energy

Bond Strength
Understanding the basis behind the periodicity of the
properties of the elements and exceptions to
periodicity.

The underlying basis for the periodicity is due to the

periodically recurring electronic structure of atoms,
which in turn causes similarities in the properties and
their correlation with atomic mass.

We need a theory and model to describe the electronic

structure about atoms in order to understand the
fundamental basis for the periodic table.

We start with the simplest theory of electronic

structure of atoms and molecules: The Lewis theory.
 Lewis Structures

1875-1946

full

valence

http://web.mit.edu/invent/iow/lewis.html
 Concept Check!
• How many valence electrons does Flourine (F) have?

1. 1
2. 2
3. 3
4. 4
5. 5
6. 6
7. 7
8. 8
9. 0
 How many valence electrons does
nitrogen (N) have?
1. 1
2. 2
3. 3
4. 4
5. 5
6. 6
7. 7
8. 8
9. 0
 The Octet Rule:

two

Bonding vs. Lone pair electrons

6 3
 Drawing Lewis Structures: e.g. HCN
1. Draw a skeleton structure, which atom goes in the middle?
– Hydrogen or Fluorine will almost always be terminal atoms.
– For other atoms, in order to draw the lowest energy Lewis structure, the atom with the
lowest ionization energy is put in the center.
– Which atom would be in the center of the Lewis structure for HCN?
– Carbon!

2. Count the total number of valence electrons. If there is a negative ion, add the absolute value of total
charge to the count of valence electrons; if positive ion, subtract.

3. Count the total # of e-s needed for each atom to have a full valence shell.

4. Subtract the number in step 2 (valence electrons) from the number in step 3(total electrons for full
shells). The result is the number of bonding electrons.

5. Assign 2 bonding electrons to each bond.

6. If bonding electrons remain, make some double or triple bonds. In general, double bonds form
only between C, N, O, and S. Triple bonds are usually restricted to C, N, and O.

7. If valence electrons remain, assign them as lone pairs, giving octets to all atoms except hydrogen.

8. Determine the formal charge. FC = V – L -(½)S

http://www.kentchemistry.com/links/bonding/lewisdotstruct.ht
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Example: H C N

2) 1 + 4 + 5 = 10 valence electrons

3) 2 + 8 (2) = 18 electrons needed for octet

4) 18 – 10 = 8 “bonding” electrons
7) 10 – 8 = 2 “lone” electrons
 Formal Charge
Formal charge is a measure of the extent to which an atom has gained or lost
an electron in the process of forming a covalent bond.

FC = V – L -(½)S

V ≡ number of valence electrons

L ≡ number of Lone Pair electrons
S ≡ number of Shared (bonding) electrons

For an electronically-neutral molecule, the sum of the formal charges of the individual
atoms must be zero.

Structures with least charge separation are the most stable structures in terms of
energy.
 Using FC to Decide Between
Different Lewis Structures
SCN-

IEC = 1090 k, IES 1000, IEN = 1400

5 - 4 – 2 = -1 -1 +1
4–0-4=0 -2 -2
6–4-2=0 +2 0

Most Stable!
 Resonance Structures
For certain molecules, more than one Lewis structure is needed to correctly describe
the valence electron structure of the molecule.

We might expect one short O=O bond and one long O-O bond, but experimental evidence
demonstrates that the two bonds are equal.

Resonance structures are two (or more) structures with the same arrangement of atoms
but a different arrangement of electrons.
 Breakdown of the Octet Rule
• Case 1. Odd number of valence electrons

For molecules with an odd number of valence electrons, it is not possible for each atom in the
molecule to have an octet, since the octet rule works by pairing e-s.

Example: CH3

2) 3(1) + 4 = 7 valence electrons

3) 3(2) + 8 = 14 electrons needed for octet

4) 14 – 7 = 7 “bonding” electrons

Radical species: molecule with an unpaired electron.

Radicals are usually very reactive. The reactivity of radical species leads to interesting (and
sometimes harmful) biological activity.
 Nitric Oxide
Some radicals are more stable. For example, NO

2) 5 + 6 = 11 valence electrons
3) 8 + 8 = 16 electrons needed for octet
4) 16 – 11 = 5 bonding electrons

http://www.biologynews.net/archives/2007/10/30/elevated_nitric_oxide_in_blood_is_key_to_high_altitude_function_for_tibetans.htm
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 Case II: Octet Deficient
Molecules
Some molecules are stable with an incomplete octet. Group 13 elements B and
Al have this property. BF 3

2) 3 + 3(7) = 24 valence electrons 5) assign two electrons per bond.

3) 8 + 3(8) = 32 electrons needed for octet 6) 8 – 6 = 2 extra bonding electrons
4) 32 – 24 = 8 bonding e-s 7) 24 – 8 = 16 lone pair electrons

8) calculate formal charges:

4 4 +1

6 2 0

experiments suggest that all three B-F bonds have the same length, that of a single bond.
 New Structure!

The formal charges are more favorable for this structure!

Case III Valence Shell Expansion
Elements with n = or > 3 have empty d-orbitals, which means more than eight electrons
can fit around the central atom.

large and is
Expanded valence shells are more common when the central atom is ______
bonded to small, highly electronegative atoms such as O, F, and Cl.

Consider PCl5

40
48
8

To make five P-Cl bonds, need ten shared electrons. So 40 – 10 = 30 lone-pair electrons.
 Chromate Ion
Consider
CrO42-
2 32

40
8

Crocoite specimen from the Red Lead Mine,

TasmaniaCrocoite specimen from the Red Lead Mine,
Tasmania, Australia
Formal Charge on the central atom should either be zero or equal to the charge
on molecule! If not the structure is not stable!

Also experimentally, Cr-O bond length and strength are between that of a single and
double bond!
 Actual Structure

four

How many more resonance structures would you expect??

Valence shell expansion around Cr results in less formal charge separation. More stable
Lewis structure.
 Identify which of the following
molecules are radical species: CO and
OH.

1. CO
2. OH
3. CO and OH
4. Neither are
radicals.
The periodic table by electronegativity

Metals: Low tendency to attract electrons,

high tendency to release electrons.
Electron Donors

Non-metals: High tendency to attract

electrons, low tendency to release
electrons. Electron Acceptors
 Polar Covalent Bond:
Perfectly-ionic and perfectly-covalent bonds are the two extremes of
bonding. In reality, most bonds fall somewhere in the middle.
A polar covalent bond is an ____________
unequal sharing of electrons between
two atoms with different electronegativities (χ).

δ is fraction of a full charge (e) that is asymmetrically distributed.

H2 is a “perfect” covalent bond, δ = 0.

 Polar Covalent Bond; Definition
• A polar covalent bond is an unequal sharing of e-s between two atoms
with with different electronegativities (χ).

• In general, a bond between two atoms with an χ difference of > 0.4 and <
1.7(on the Pauling scale) is considered polar covalent.
Dipole Moment
Examples of Polar Molecules:
• Polar molecules have a non-zero net dipole moment.

non-polar
polar
 Polar Groups and Solubility

B-9

Folic acid
fat water
 Nitrous Oxide
(-1)

(-1)
(+1) (+1)

The nitrogen-nitrogen bond is 1.126 Angstroms long which is slightly

longer than the triple bond length in N2 (1.098 Angstroms).

The nitrogen-oxygen bond is 1.186 Angstroms long. This is longer than the
typical N=O bond (about 1.14 Angstroms), which agrees with the prediction
of partial single-bond character for the NO bond in N2O.

Both structures are weakly polar but in opposite directions, with a negative FC
being on Nitrogen on one structure and Oxygen on the other structure.

The low polarity of the gas makes it both oil and water soluble! This allows it to
travel through the bloodstream and into the fatty membranes of nerve cells where
it produces its characteristic effects.

http://antoine.frostburg.edu/chem/senese/101/inorganic/faq/print-laughing-gas.shtm
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 Why does NOS make the car
go faster?
• Heating nitrous oxide to about 570 degree F, it splits into oxygen and
nitrogen.
• The injection of nitrous oxide into an engine therefore means more oxygen is
available for combustion
• More oxygen, means you can also inject more gasoline, and that means
more energy, power and speed.
• WARNING Nitrous fitted to smaller engines without modified pistons can
blow the engine to pieces!

power from 0.5 hp – 3000 hp!